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In Atomic Physics , the electron configuration is the arrangement of Electron s in an Atom , Molecule or other body. Specifically, it is the placement of electrons into Atomic , Molecular , or other forms of Electron Orbital s. ELECTRON CONFIGURATION IN ATOMS ''The discussion below presumes knowledge of material contained at Atomic Orbital .'' Summary of the quantum numbers The state of an electron in an atom is given by four Quantum Number s. Three of these are integers and are properties of the Atomic Orbital in which it sits (a more thorough explanation is given in that article). No two electrons in one atom can have the same set of these four quantum numbers ( Pauli Exclusion Principle ). Shells and subshells Shells and subshells (also called energy levels and sublevels) are defined by the quantum numbers, NOT by the distance of its electrons from the nucleus. In large atoms, shells above the second shell overlap (see Aufbau Principle ). States with the same value of ''n'' are related, and said to lie within the same '' Electron Shell ''. States with the same value of ''n'' and also ''l'' are said to lie within the same '' Electron Subshell ''. If the states also share the same value of ''m'', they are said to lie in the same Atomic Orbital . Because electrons have only two possible spin states, an atomic orbital cannot contain more than two electrons ( Pauli Exclusion Principle ). A subshell can contain up to 4''l''+2 electrons; a shell can contain up to 2''n''&2 electrons. Worked example Here is the electron configuration for a filled fifth shell: This information can be written as 5s2 5p6 5d10 5f14 5g18 (see below for more details on notation). The subshell labels s, p, d, and f originate from a now-discredited system of categorizing Spectral Lines as "sharp", "principal", "diffuse", or "fundamental", based on their observed Fine Structure . When the first four types of orbitals were described, they were associated with these spectral line types, but there were no other names. The designation ''g'' was derived by following alphabetical order. Shells with more than five subshells are theoretically permissible, but this covers all discovered elements. Notation Physicists and chemists use a standard notation to describe atomic electron configurations. In this notation, a subshell is written in the form nxe, where n is the shell number, x is the subshell label and e is the number of electrons in the subshell. An atom's subshells are written in order of increasing energy - in other words, the sequence in which they are filled (see Aufbau principle below). For instance, ground-state Hydrogen has one electron in the s subshell of the first shell, so its configuration is written 1s1. Lithium has two electrons in the 1s subshell and one in the (higher-energy) 2s subshell, so its ground-state configuration is written 1s2 2s1. Phosphorus (atomic number 15), is as follows: 1s2 2s2 2p6 3s2 3p3. For atoms with many electrons, this notation can become lengthy. It is often abbreviated by noting that the first few subshells are identical to those of one or another Noble Gas . Phosphorus, for instance, differs from Neon (1s2 2s2 2p6) only by the presence of a third shell. Thus, the electron configuration of neon is pulled out, and phosphorus is written as follows: {Link without Title} 3s2 3p3. An even simpler version is simply to quote the number of electrons in each shell, e.g. (again for phosphorus): 2-8-5. Aufbau principle In the Ground State of an atom (the condition in which it is ordinarily found), the electron configuration generally follows Aufbau Principle . According to this principle, electrons enter into states in order of the states' increasing energy; i.e., the first electron goes into the lowest-energy state, the second into the next lowest, and so on. The order in which the states are filled is as follows: A pair of electrons with identical spins has slightly more energy than a pair of electrons with opposite spins. Since two electrons in the same orbital must have opposite spins, this causes electrons to prefer to occupy different orbitals. This preference manifests itself if a subshell with (one that contains more than one orbital) is less than full. For instance, if a ''p'' subshell contains four electrons, two electrons will be forced to occupy one orbital, but the other two electrons will occupy ''both'' of the other orbitals, and their spins will be equal. This phenomenon is called '' Hund's Rule ''. The Aufbau principle can be applied, in a modified form, to the Proton s and Neutron s in the Atomic Nucleus (see the Shell Model of Nuclear Physics ). Exceptions in 3d, 4d, 5d A ''d'' subshell that is half-filled or full (ie 5 or 10 electrons) is more stable than the ''s'' subshell of the next shell. This is the case because it takes less energy to maintain an electron in a half-filled ''d'' subshell than a filled ''s'' subshell. For instance, Copper (atomic number 29) has a configuration of 3d10, not [Ar 4s2 3d9 as one would expect by the Aufbau principle. Likewise, Chromium (atomic number 24) has a configuration of 3d5, not [Ar 4s2 3d4. This can be most easily understood by stepping through the electron configuration shown at {Link without Title} Period 5th has more exceptions: This can be seen by stepping through the electron configuration shown at {Link without Title} This can be seen by stepping through the electron configuration shown at {Link without Title} Relation to the structure of the periodic table See Also: Periodic table#Explanation of the structure of the periodic table Electron configuration is intimately related to the structure of the Periodic Table . The chemical properties of an atom are largely determined by the arrangement of the electrons in its outermost ("valence") Shell (although other factors, such as Atomic Radius , Atomic Mass , and increased accessibility of additional electronic states also contribute to the chemistry of the elements as atomic size increases). ELECTRON CONFIGURATION IN MOLECULES In molecules, the situation becomes more complex, as each molecule has a different orbital structure. See the Molecular Orbital article and the Linear Combination Of Atomic Orbitals method for an introduction and the Computational Chemistry article for more advanced discussions. ELECTRON CONFIGURATION IN SOLIDS In a Solid , the electron states become very numerous. They cease to be discrete, and effectively blend together into continuous ranges of possible states (an Electron Band ). The notion of electron configuration ceases to be relevant, and yields to Band Theory . SEE ALSO
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